General Characteristics of Alkaline Earth Metals

Physical state and occurrence. Beryllium is a rare element, and radium is the rarest of all the group-2 elements, occurring only in minute amounts (chiefly as a decay product wherever uranium-bearing ores are found, since radium is itself a decay product of uranium). Magnesium and calcium, by contrast, are very common in the earth's crust -- calcium is the fifth most abundant element overall, and magnesium the eighth. Both are found in many rocks and minerals: magnesium in carnallite, magnesite and dolomite; calcium in chalk, limestone and gypsum. Most strontium occurs in the minerals celestite and strontianite, while barium is somewhat less common, occurring mostly as the mineral barite (Table 5.7).

A note on fireworks: several alkali and alkaline earth metals are responsible for firework colours. Combined with chlorine, barium gives a green spark, calcium gives orange, and lithium gives a medium red; strontium carbonate gives a bright red, sodium nitrate gives orange, potassium/rubidium nitrates give violet, and caesium gives an indigo colour. In every case, the heat of burning "excites" the metal's electrons to a higher-than-normal energy level, and the colourful burst of light comes from the extra energy released as they fall back. Blue is the hardest firework colour to produce, since the compound responsible (copper chloride) breaks down in a hot flame; pyrotechnicians have used magnalium (a magnesium-aluminium alloy) to brighten firework colours generally, though a blue as vivid as the reds, greens and yellows remains elusive.

Electronic configuration. Every alkaline earth metal has two electrons in the valence shell, preceded by a noble-gas core, giving the general configuration [Noble gas]ns², where n is the valence shell (Table 5.8).

Atomic and ionic radii. The atomic and ionic radii of the alkaline earth metals are smaller than those of the corresponding alkali metals in the same period. This is because group 2 elements have a higher nuclear charge, which pulls the electrons in more strongly. Down the group, however, the radii still increase, due to the gradual addition of extra shells and the resulting increase in the screening effect (Table 5.9).

Common oxidation state. With two electrons in the valence shell, group 2 elements achieve the stable noble-gas configuration by losing both of them -- so every alkaline earth metal shows a single oxidation state, +2, in its compounds.

Ionisation enthalpy. Because of their fairly large atomic size, alkaline earth metals have low ionisation enthalpies compared to the p-block elements, and (as with the alkali metals) the ionisation enthalpy falls down the group as atomic size increases -- the result of both an added shell and a greater screening effect from the inner-shell electrons. Compared with the alkali metals of the same period, though, group 2 members have higher ionisation-enthalpy values, since their smaller size pulls the valence electrons in more strongly -- correspondingly, alkaline earth metals are less electropositive than the alkali metals (Figure 5.6 plots the smooth fall in first ionisation energy from Be to Ba).

Interestingly, while the first ionisation enthalpy of an alkaline earth metal is higher than that of the corresponding alkali metal, its second ionisation enthalpy is much lower than the alkali metal's second ionisation enthalpy. This is because, for an alkali metal, the second electron has to be pulled out of a cation that has already achieved a stable noble-gas configuration -- a very costly step. For an alkaline earth metal, by contrast, the second electron is removed from a monovalent cation that still has one electron left in its outermost shell, which is comparatively easy to remove. …